Question

(a) According to the Bohr model, why might someone expect that atoms would not change in size as you go from left to right across a period? (b) In fact, atoms do change in size across a period. What is the trend and what is the explanation for it?

Short answer

(a) According to the Bohr model, one might expect atoms not to change in size across a period because electrons are added to the same shell and the shell number remains constant. The Bohr model doesn't consider shielding effect and electron-electron repulsion. (b) However, atomic size does change across a period, with the trend being a decrease in atomic radius from left to right. This occurs due to increasing nuclear charge, which attracts electrons in the same shell closer to the nucleus, overcoming the shielding effect and electron-electron repulsion, thereby making atoms smaller across the period.

Step-by-step solution

Part (a): Reason for expecting no change in atomic size according to Bohr model

According to the Bohr model, electrons are organized into concentric shells around the nucleus. As we go across a period (left to right) in the periodic table, electrons are added to the same shell (same principal quantum number) and the number of protons in the nucleus also increases by one unit. Since the shell number remains the same, one might expect that the electron cloud and, consequently, the atom's size should remain constant across a period. Moreover, the Bohr model does not take into account the shielding effect and electron-electron repulsion, which are important factors affecting the atomic size.

Part (b): Actual trend in atomic size across a period and explanation

In reality, the atomic radius decreases as we go from left to right across a period in the periodic table. This is due to the increase in the number of protons in the nucleus, which increases the effective nuclear charge attracting the electrons in the same shell closer to the nucleus. The shielding effect by inner electrons is less significant than the effect of increased nuclear charge since electrons within the same shell don't shield each other very effectively. Due to these factors, electron clouds are pulled inwards, making the atoms smaller as we move across a period.

Key concepts

Bohr Model

The Bohr model is a classic concept in atomic theory that describes the atom as a miniature solar system. In this model, electrons orbit around the nucleus in defined paths or shells. These shells are like concentric circles around the nucleus, where each electron has a specific energy level.

What makes the Bohr model notable is that it allows for easy visualization of electron configuration. According to Bohr, as you move from left to right across a period on the periodic table, electrons are added to these same energy levels (or shells). This might seem like the atom's size should remain constant because the electrons occupy the same shell.

However, Bohr's model doesn't account for the entire story. It overlooks interactions like shielding and effective nuclear charge that also play roles in determining atomic size. Despite its simplifications, the Bohr model was revolutionary in setting the groundwork for understanding atomic structures.

Atomic Size

Atomic size, often referred to as atomic radius, is essentially the size of an atom. It’s measured from the center of the nucleus to the boundary of the electron cloud. Atomic size can vary depending on the components of the atom, such as the number of shells and the number of electrons.

The trend across a period shows a decrease in atomic size from left to right. This reduction is not intuitive if you only consider that electrons fill the same shell. However, as more protons are added to the nucleus while moving across a period, the pull on the electron cloud increases, pulling electrons closer to the nucleus and thereby decreasing the size of the atom.

In conclusion, the atomic size becomes smaller across a period due to the increase in nuclear charge, creating a stronger attraction between the nucleus and the electron cloud.

Effective Nuclear Charge

Effective nuclear charge ( Z_{eff} ) is an important concept to understand the behavior of electrons in an atom. It is the net positive charge experienced by an electron in a multi-electron atom. Simply put, it accounts for the positive charge from the nucleus that an electron "feels" after considering shielding effects.

As you go from left to right across a period, the effective nuclear charge increases. This is because the number of protons in the nucleus increases, enhancing the attractive force exerted on the electrons in the same energy level. Despite more electrons being present, these new additions are not efficient at shielding each other from the increased positive charge of the nucleus.

Thus, the increased effective nuclear charge across a period contributes to the decrease in atomic size by pulling the electron cloud closer to the nucleus. This explains why atoms do not stay the same size across a period as one might expect from a simple shell model.

Shielding Effect

The shielding effect describes how inner electrons can obscure outer electrons from the full force of the nuclear charge. It acts as a barrier between the nucleus and the valence shell electrons, reducing the nucleus’s pull on those outer electrons.

This effect is more pronounced when inner electron shells shield valence electrons, allowing them to reside at greater distances from the nucleus.

However, within the same energy level, electrons don't shield each other effectively. That's why even as electrons are added across a period, they don't significantly reduce the increasing pull of the nucleus. The net result is that despite more electrons added to the existing shells, these electrons are drawn closer to the nucleus due to increased effective nuclear charge, leading to a reduction in atomic size across a period.

Understanding the shielding effect is crucial for grasping why atomic size changes in surprising ways depending on where you look in the periodic table.